Meaning of CHEMICAL REACTION in English

CHEMICAL REACTION

any type of chemical process in which substances are changed into different substances, as differentiated from other kinds of changes-those of position or of form-undergone by matter. Chemical reactions are manifested by the disappearance of properties characteristic of the starting materials and the appearance of new properties that distinguish the products; within the limits of observation, the mass of the products formed is equal to the mass of the substances consumed. Chemical reactions were no doubt known to prehistoric cultures, and speculations on the causes of such changes in the material world led certain ancient Greek philosophers to suppose that matter is perhaps composed of discrete but minute particles. This so-called atomic theory was revived in the early 19th century in England by John Dalton, who postulated that atoms-minute, indivisible, and unalterable particles-are the smallest units of matter, that the properties of a substance depend on the types of atoms present in it, and that all atoms of a particular element are identical. Dalton then argued that a chemical reaction was the redistribution of atoms-that is, the dissociation of atoms initially conjoined and the rejoining of those atoms in a new arrangement. During this same period the French chemist and physicist Joseph-Louis Gay-Lussac, among others, discovered that the ratio of the volumes of reacting gases could always be expressed in small whole numbers. In 1811 Amadeo Avogadro of Italy suggested that this relationship could be explained by assuming that gases of equal volume, temperature, and pressure contain the same number of particles. Initially ignored, Avogadro's theory came to be accepted by 1860, largely through the efforts of another Italian chemist, Stanislao Cannizzaro. Later research by chemists and physicists revealed that atoms are not indivisible but made of smaller particles of opposite electrical charge and vastly different masses. The smaller, lighter charged particle is now called the electron, and the heavier particle of opposite charge is called the nucleus. It is now known that the dissociation and regrouping of atoms or groups of atoms are mediated by the interaction of their electrons. This interaction may be expressed by sharing electrons between or among atoms or groups of atoms in a molecule, or it may be expressed by transferring electrons from one atom or one group of atoms to another. The former interaction is called a covalent bond; the latter, ionic. A chemical reaction, then, involves the absorption of energy by the bond-forming electrons, causing the bonds to rupture. Under these conditions opportunities are created for new bonds to form; when they do, energy is released. If in a chemical reaction the energy needed to break a bond is less than that released by the formation of a new bond, the reaction is termed exoergic (or exothermic). The converse is true in endoergic (or endothermic) reactions. In rare instances, termed aergic (or athermic) reactions, there is no net absorption or evolution of energy. In general, some form of energy must be supplied to reactants to initiate bond rupture. This energy, called activation energy, can be in the form of heat, electromagnetic radiation, or electrical energy. When allowed to proceed in the presence of certain substances, chemical bonds can be ruptured with less activation energy than normally required. Such substances, which are not consumed during the reaction, are called catalysts; the reactions that they facilitate are said to be catalyzed. Although the speed of a reaction is related to the energy relationships discussed above, whether a reaction will occur at all depends on changes in the entropy both of the reactants and products, and of their surroundings. Entropy is a measure of the energy unavailable to do work in a system; i.e., the energy associated with disorder. A chemical reaction cannot occur unless there is an overall increase in entropy. any type of chemical process in which substances are changed into different substances, as differentiated from other kinds of changes-those of position or of form-undergone by matter. Chemical reactions are manifested by the disappearance of properties characteristic of the starting materials and the appearance of new properties that distinguish the products; within the limits of observation, the mass of the products formed is equal to the mass of the substances consumed. Thus, when wood burns, the substances present initially, wood and oxygen in the atmosphere, are converted in a chemical reaction to water vapour, carbon dioxide, and ash. All combustions are chemical reactions. Other types of familiar chemical reactions include decay, fermentation, the hardening of cement, the development of a latent image in an exposed photographic film, the tarnishing of silver, the corrosion of steel, the evolution of gas when vinegar and soda are mixed, the synthesis of nylon, and the digestion of food. In a general sense, material substances can undergo change in three ways: a change of position, called movement; a change of form, such as the freezing of liquid water; and a change of substance, a chemical reaction. Some classify changes of form as chemical reactions, but, historically, the term chemical reaction has been applied only to changes of substance. The application to change of form is discussed below. Using the historical definition, each different chemical reaction displays the same unique characteristics. Additional reading General works Overviews are provided by William F. Kieffer, Chemistry: A Cultural Approach (1971), an introductory text for the interested person who is uninstructed in science; Harold G. Cassidy, Science Restated: Physics and Chemistry for the Non-Scientist (1970), a philosophical approach, with emphasis on the past and future contributions of physics and chemistry to culture; Cooper H. Langford and Ralph A. Beebe, The Development of Chemical Principles (1969), outstanding in clarity and rigour, but more advanced than Kieffer or Cassidy; J. Arthur Campbell, Why Do Chemical Reactions Occur? (1965), a well-written exposition; Edward L. King, How Chemical Reactions Occur (1963), an elementary treatment of kinetics and mechanisms, chain reactions, activation energy, and the use of laboratory instruments; and Roman Mierzecki, The Historical Development of Chemical Concepts, trans. from Polish (1991). A historical approach that considers the influence of Linus Pauling is found in Ahmed Zewail (ed.), The Chemical Bond: Structure and Dynamics (1992). Introductory college-level textbooks in general and organic chemistry provide much useful information on chemical reactions; these include Henry F. Holtzclaw, Jr., and William R. Robinson, General Chemistry, 8th ed. (1988); T.R. Dickson, Introduction to Chemistry, 6th ed. (1991); Ralph J. Fessenden and Joan S. Fessenden, Organic Chemistry, 5th ed. (1994); Robert Thornton Morrison and Robert Nielson Boyd, Organic Chemistry, 6th ed. (1992); and Andrew Streitwieser, Jr., Introduction to Organic Chemistry, 4th ed. (1992). An advanced-level text that also gives bibliographic references is R. Stephen Berry, Stuart A. Rice, and John Ross, Physical Chemistry (1980), especially part 3, which is a thorough exposition of chemical reactions in gases and solutions, with a grounding in the collision processes that lead to reactions. Jay A. Young The Editors of the Encyclopdia Britannica Mechanisms and kinetics Books of general interest include C.K. Ingold, Structure and Mechanism in Organic Chemistry, 2nd ed. (1969); Fred Basolo and Ralph G. Pearson, Mechanisms of Inorganic Reactions, 2nd ed. (1967); Ronald Breslow, Organic Reaction Mechanisms, 2nd ed. (1969); Calvin D. Ritchie, Physical Organic Chemistry, 2nd ed. rev. and expanded (1990); Robert B. Jordan, Reaction Mechanisms of Inorganic and Organometallic Systems (1991); Dimitris Katakis and Gilbert Gordon, Mechanisms of Inorganic Reactions (1987); P.J. Robinson and K.A. Holbrook, Unimolecular Reactions (1972), a thorough discussion of how energy gets shuffled within a molecule to lead to reaction; Zdenek Slanina, Contemporary Theory of Chemical Isomerism (1986), a broad survey of rearrangement mechanisms; Cooper H. Langford and Harry B. Gray, Ligand Substitution Processes (1966); Andrew Streitwieser, Jr., Molecular Orbital Theory for Organic Chemists (1961, reissued 1974); T.L. Gilchrist and R.C. Storr, Organic Reactions and Orbital Symmetry, 2nd ed. (1979); Ralph G. Pearson, Symmetry Rules for Chemical Reactions (1976); J.N. Murrell et al., Molecular Potential Energy Functions (1984), a survey of methods and results; D. Michael P. Mingos and David J. Wales, Introduction to Cluster Chemistry (1990), reactions and mechanisms in atomic and molecular clusters; John W. Moore and Ralph G. Pearson, Kinetics and Mechanism, 3rd ed. (1981); James H. Espenson, Chemical Kinetics and Reactions Mechanisms (1981); Frank Wilkinson, Chemical Kinetics and Reaction Mechanisms (1980); Keith J. Laidler, Chemical Kinetics, 3rd ed. (1987); Sidney W. Benson, The Foundations of Chemical Kinetics (1960, reprinted 1982); R.P. Bell, The Proton in Chemistry, 2nd ed. (1973); E. Buncel and C.C. Lee (eds.), Isotopic Effects: Recent Developments in Theory and Experiment (1984); John E. Leffler and Ernest Grunwald, Rates and Equilibria of Organic Reactions as Treated by Statistical, Thermodynamic, and Extrathermodynamic Methods (1963, reprinted 1989); and Barry K. Carpenter, Determination of Organic Reaction Mechanisms (1984).Books of specialized interest include C.A. Bunton, Nucleophilic Substitution at a Saturated Carbon Atom (1963); D.V. Banthorpe, Elimination Reactions (1963); D. Bethell and V. Gold, Carbonium Ions (1967); S.R. Hartshorn, Aliphatic Nucleophilic Substitution (1973); J. Milton Harris and Samuel P. McManus (eds.), Nucleophilicity (1987); Pierre Vogel, Carbocation Chemistry (1985); Herbert C. Brown and Paul von R. Schleyer, The Nonclassical Ion Problem (1977); Robert B. Bates and Craig A. Ogle, Carbanion Chemistry (1983); E. Buncel and T. Durst (eds.), Comprehensive Carbanion Chemistry (1980- ); George A. Olah, Ripudaman Malhotra, and Subhash C. Narang, Nitration: Methods and Mechanisms (1989); Kenneth Schofield, Aromatic Nitration (1980); R. Taylor, Electrophilic Aromatic Substitution (1990); and Peter B.D. De La Mare, Electrophilic Halogenation (1976).Current information may be found in the following series: Progress in Physical Organic Chemistry (irregular); Advances in Physical Organic Chemistry (annual); Topics in Stereochemistry (irregular); Advances in Carbocation Chemistry (annual); Advances in Carbanion Chemistry (irregular); and Advances in Detailed Reaction Mechanisms (annual); and in reviews on various topics in Annual Reports on the Progress of Chemistry; and Chemical Reviews (bimonthly).Works on chemical kinetics include the text by Moore and Pearson, cited above; Samuel Glasstone, Keith J. Laidler, and Henry Eyring, The Theory of Rate Processes (1941), the earliest comprehensive treatment of the theory of absolute reaction rates; and Henry Eyring, S.H. Lin, and S.M. Lin, Basic Chemical Kinetics (1980). Explorations of chemical relaxation phenomena include Discussions of the Faraday Society, no. 17 (1954), a colloquium on the study of fast reactions, which includes a discussion of chemical relaxation by Manfred Eigen; Molecular Relaxation Processes (1966), papers from a Chemical Society symposium emphasizing the use of relaxation to determine molecular structures; M. Eigen and L. De Mayer, "Theoretical Basis of Relaxation Spectroscopy," Techniques of Chemistry, vol. 6 (1973), with other chapters discussing specific techniques involving relaxation methods; Francis K. Fong, Theory of Molecular Relaxation (1975); J. Fnfschilling (ed.), Relaxation Processes in Molecular Excited States (1989); and D. Steele and J. Yarwood (eds.), Spectroscopy and Relaxation of Molecular Liquids (1991).Eric K. Rideal and Hugh S. Taylor, Catalysis in Theory and Practice, 2nd ed. (1926), recounts the state of this subject during World War I and the rapid advances in theory that occurred shortly thereafter. Paul H. Emmett, Catalysis Then and Now: A Survey of the Advances in Catalysis (1965), is a classic work written by a Nobel Prize winner. Paul H. Emmett (ed.), Catalysis, 7 vol. (1954-60), is a monumental treatment of the subject, both theory and practice, contributed by a group of international experts in the field. Additional information may be found in William P. Jencks, Catalysis in Chemistry and Enzymology (1969, reissued 1987); R. Pearce and W.R. Patterson (eds.), Catalysis and Chemical Processes (1981); John M. Thomas and Kirill I. Zamaraev (eds.), Perspectives in Catalysis (1992); Bruce C. Gates, Catalytic Chemistry (1992); and J.A. Moulijn, P.W.N.M. van Leeuwen, and R.A. van Santen (eds.), Catalysis: An Integrated Approach to Homogeneous, Heterogeneous, and Industrial Catalysis (1993). Advances in Catalysis (annual) records the year-to-year advances in the field as reported by international authorities. Peter B.D. de la Mare The Editors of the Encyclopdia Britannica Henry Eyring The Editors of the Encyclopdia Britannica Larry D. Faller The Editors of the Encyclopdia Britannica Sir Hugh S. Taylor The Editors of the Encyclopdia Britannica Mechanisms and kinetics Chemical relaxation phenomena The term relaxation is used by chemists and physicists to describe the interval or time lag between the application of an external stress to a system-that is, to an aggregation of matter-and its response. The relaxation effect may be caused by a redistribution of energy among the nuclear, electronic, vibrational, and rotational energy states of the atoms and molecules that comprise the system, or it may result from a shift in the ratio of the number of product molecules to the number of reactant molecules (those initially taking part) in a chemical reaction. The measurement of relaxation times can provide many insights into atomic and molecular structures and into the rates and mechanisms of chemical reactions. Historical survey The word relaxation was originally applied to a molecular process by the English physicist James Clerk Maxwell. In a paper "On the Dynamical Theory of Gases," which he presented in 1866, Maxwell referred to the time required for the elastic force produced when fluids are distorted to diminish or decay to 1/e (e is the base of the natural logarithm system) times its initial value as the "time of relaxation" of the elastic force. The earliest suggestion of a chemical relaxation effect is contained in a dissertation (Berlin, 1910) based on research directed by the German physical chemist Walther Nernst. Measurements of sound propagation through the gas nitrogen tetroxide, which breaks up, or dissociates, into nitrogen dioxide, led Nernst to suggest that experiments at frequencies at which the dissociation reaction could not keep pace with the temperature and pressure variations that occur within a sound wave, would permit evaluation of the dissociation rate. Ten years later, at a meeting of the Prussian Academy of Sciences, Albert Einstein presented a paper in which he described the various theoretical aspects of this relaxation effect. The detection of the chemical relaxation effect predicted by Nernst and Einstein did not become technically feasible until the last half of the 20th century. In the first half of the century physicists and chemists in studying relaxation concentrated on physical relaxation processes. Peter Debye referred to the time required for dipolar molecules (ones whose charges are unevenly distributed) to orient themselves in an alternating electric field as dielectric relaxation. Sound absorption by gases was used to investigate energy transfer from translational, or displacement in space, to rotational (spinning and tumbling) and vibrational (oscillations within the molecule) degrees of freedom, the three independent forms of motion for a molecule. The former requires only a few molecular collisions, but the transfer of energy between translational and vibrational modes may require thousands of collisions. Because the processes are not instantaneous but time dependent, relaxation effects are observed. Their measurement provides information about molecular bonding and structure. Chemical relaxation was rediscovered by the German physical chemist Manfred Eigen in 1954. Since then, technological advances have permitted the development of techniques for the measurement of relaxation times covering the entire range of molecular processes and chemical reactivity. The great variety of relaxation phenomena and of the techniques developed for their study precludes a comprehensive survey. To facilitate a general discussion, the relaxing system, its initial and final states, the nature of the disturbance, and the system's response are considered separately. Examples are cited that emphasize the important features of relaxation phenomena and illustrate the variety of information that can be obtained from their study. A moderately detailed description of one relaxation technique, the temperature-jump method, is used to summarize the discussion. Mechanisms and kinetics Catalysis The rates of chemical reactions-that is, the velocities at which they occur-depend upon a number of factors, including the chemical nature of the reacting species and the external conditions to which they are exposed. A particular phenomenon associated with the rates of chemical reactions that is of great theoretical and practical interest is catalysis. As noted earlier, this phenomenon is the acceleration of chemical reactions by substances not consumed in the reactions themselves-substances known as catalysts. The study of catalysis is of interest theoretically because of what it reveals about the fundamental nature of chemical reactions; in practice, the study of catalysis is important because many industrial processes depend upon catalysts for their success. Finally, the peculiar phenomenon of life would hardly be possible without the biological catalysts termed enzymes. History The term catalysis (from the Greek kata-, "down," and lyein, "loosen") was first employed by the great Swedish chemist Jns Jacob Berzelius in 1835 to correlate a group of observations made by other chemists in the late 18th and early 19th centuries. These included: the enhanced conversion of starch to sugar by acids first observed by Gottlieb Sigismund Constantin Kirchhoff; Sir Humphry Davy's observations that platinum hastens the combustion of a variety of gases; the discovery of the stability of hydrogen peroxide in acid solution but its decomposition in the presence of alkali and such metals as manganese, silver, platinum, and gold; and the observation that the oxidation of alcohol to acetic acid is accomplished in the presence of finely divided platinum. The agents promoting these various reactions were termed catalysts, and Berzelius postulated a special, unknown catalytic force operating in such processes. In 1834 the English scientist Michael Faraday had examined the power of a platinum plate to accomplish the recombination of gaseous hydrogen and oxygen, the products of electrolysis of water, and the retardation of that recombination by the presence of other gases, such as ethylene and carbon monoxide. Faraday maintained that essential for activity was a perfectly clean metallic surface (at which the retarding gases could compete with the reacting gases and so suppress activity), a concept that would later be shown to be generally important in catalysis. Many of the primitive technical arts involved unconscious applications of catalysis. The fermentation of wine to acetic acid, the manufacture of soap from fats and alkalies, and the formation of ether from alcohol and sulfuric acid-all catalytic reactions-were well known in man's early history. Sulfuric acid prepared by firing mixtures of sulfur and nitre (sodium nitrate) was an early forerunner of the lead chamber process of sulfuric acid manufacture, in which sulfur dioxide oxidation was accelerated by the addition of oxides of nitrogen. (A mechanism for the latter process was suggested by Sir Humphry Davy in 1812 on the basis of experiments carried out by others.) In 1850 the concept of a velocity of reaction was developed during studies of hydrolysis, or inversion, of cane sugar. The term inversion refers to the change in rotation undergone by monochromatic light when it is passed through the reaction system, a parameter that is easily measured, facilitating study of the reaction. It was found that the rate of inversion was, at any moment, proportional to the amount of cane sugar undergoing transformation and that the rate was accelerated by the presence of acids. (Later it was shown that the rate of inversion was directly proportional to the strength of the acid.) This work was, in part, the precursor of later studies of reaction velocity and the accelerating influence of higher temperature on that velocity by J.H. van't Hoff, Svante Arrhenius, and Wilhelm Ostwald, all of whom played leading roles in the developing science of physical chemistry. Ostwald's work on reaction velocities led him in the 1890s to define catalysts as substances that change the velocity of a given chemical reaction without modification of the energy factors of the reaction. This statement of Ostwald was a memorable advance since it implied that catalysts do not change the position of equilibrium in a reaction. In 1877 Georges Lemoine had shown that the decomposition of hydriodic acid to hydrogen and iodine reached the same equilibrium point at 350 C, 19 percent, whether the reaction was carried out rapidly in the presence of platinum sponge or slowly in the gas phase. This observation has an important consequence: a catalyst for the forward process in a reaction is also a catalyst for the reverse reaction. P.E.M. Berthelot, the distinguished French chemist, had confirmed this observation in 1879 with liquid systems when he found that the reaction of organic acids and alcohols, called esterification, is catalyzed by the presence of small amounts of a strong inorganic acid, just as is the reverse process-the hydrolysis of esters (the reaction between an ester and water). The application of catalysts to industrial processes was undertaken deliberately in the 19th century. P. Phillips, an English chemist, patented the use of platinum to oxidize sulfur dioxide to sulfur trioxide with air. His process was employed for a time but was abandoned due to loss of activity by the platinum catalyst. Subsequently poisons in the reactants were found to be responsible, and the process became a technical success at the turn of the 20th century. In 1871 an industrial process was developed for the oxidation of hydrochloric acid to chlorine in the presence of cupric salts impregnated in clay brick. The chlorine obtained was employed in the manufacture of bleaching powder (a dry substance that releases chlorine on treatment with acid) by reaction with lime. Again, in this reaction, it was observed that the same equilibrium was reached in both directions. Furthermore, it was found that the lower the temperature, the greater the equilibrium content of chlorine; a working temperature of 450 C produced the maximum amount of chlorine in a convenient time. Toward the close of the 19th century, the classical studies of the eminent French chemist Paul Sabatier on the interaction of hydrogen with a wide variety of organic compounds were carried out using various metal catalysts; this research led to the development of a German patent for the hydrogenation of liquid unsaturated fats to solid saturated fats with nickel catalysts. The development of three important German catalytic processes had great impact on industry at the end of the 19th century and in the early decades of the 20th. One was the so-called contact process for producing sulfuric acid catalytically from the sulfur dioxide produced by smelting operations. Another was the catalytic method for the synthetic production of the valuable dyestuff indigo. The third was the catalytic combination of nitrogen and hydrogen for the production of ammonia-the Haber process for nitrogen fixation-developed by the chemist Fritz Haber. Mechanisms and kinetics The mechanisms of chemical reactions are the detailed processes by which chemical substances are transformed into other substances. The reactions themselves may involve the interactions of atoms, molecules, ions, electrons, and free radicals, and they may take place in gases, liquids, or solids-or at interfaces between any of these. The study of the detailed processes of reaction mechanisms is important for many reasons, including the help it gives in understanding and controlling chemical reactions. Many reactions of great commercial importance can proceed by more than one reaction path; knowledge of the reaction mechanisms involved may make it possible to choose reaction conditions favouring one path over another, thereby giving maximum amounts of desired products and minimum amounts of undesired products. Furthermore, on the basis of reaction mechanisms, it is sometimes possible to find correlations among systems not otherwise obviously related. The ability to draw such analogies frequently makes it possible to predict the course of untried reactions. Finally, detailed information about reaction mechanisms permits unification and understanding of large bodies of otherwise unrelated phenomena, a matter of great importance in the theory and practice of chemistry. Generally, the chemical reactions whose mechanisms are of interest to chemists are those that occur in solution and involve the breaking and reforming of covalent bonds between atoms-covalent bonds being those in which electrons are shared between atoms. Interest in these reactions is especially great because they are the reactions by which such materials as plastics, dyes, synthetic fibres, and medicinal agents are prepared, and because most of the biochemical reactions of living systems are of this type. In addition, reactions of this kind generally occur in time scales convenient for study, neither too fast nor too slow, and under conditions that are easily manipulated for experimental purposes. Lastly, there are a number of techniques by which the mechanisms of such reactions can be investigated. Chemical reactions involve changes in bonding patterns of molecules-that is, changes in the relative positions of atoms in and among molecules, as well as shifts in the electrons that hold the atoms together in chemical bonds. Reaction mechanisms, therefore, must include descriptions of these movements with regard to spatial change and also with regard to time. The overall route of change is called the course of the reaction, and the detailed process by which the change occurs is referred to as the reaction path or pathway. Also important to the study of reaction mechanisms are the energy requirements of the reactions. Most reactions of mechanistic interest are activated processes-that is, processes that must have a supply of energy before they can occur. The energy is consumed in carrying the starting material of the reaction over an energy barrier. This process occurs when the starting material absorbs energy and is converted to an activated complex or transition state. The activated complex then proceeds to furnish the product of the reaction without further input of energy-often, in fact, with a release of energy. Such considerations are important to an understanding of reaction mechanisms because the actual course that any reaction follows is the one that requires the least energy of activation. This reaction course is not always the one that would seem simplest to the chemist without detailed study of the different possible mechanisms. The study of reaction mechanisms is complicated by the reversibility of most reactions (the tendency of the reaction products to revert to the starting materials) and by the existence of competing reactions (reactions that convert the starting material to other than the desired products). Another complicating factor is the fact that many reactions occur in stages in which intermediate products (intermediates) are formed and then converted by further reactions to the final products. In examining chemical reactions it is useful to consider several general subjects: (1) factors that influence the course of chemical reactions, (2) energy changes involved in the course of a typical reaction, (3) factors that reveal the mechanism of a reaction, and (4) the classification of reaction mechanisms. With this information in mind it is then possible to look briefly at some of the more important classes of reaction mechanisms. (The articles acid-base reaction, oxidation-reduction reaction, and electrochemical reaction deal with the mechanisms of reactions not described in this article.) General considerations Determinants of the course of reaction The reactants In analyzing the mechanism of a reaction, account must be taken of all the factors that influence its course. After the bulk chemical constituents have been identified by ordinary methods of structure-determination and analysis, any prereaction changes involving the reactants, either individually or together, must be investigated. Thus, in the cleavage of the substance ethyl acetate by water (hydrolysis), the actual reagent that attacks the ethyl acetate molecule may be the water molecule itself, or it may be the hydroxide ion (OH-) produced from it. The hydrolysis of ethyl acetate can be represented by the following equation: in which the structures of the molecules are represented schematically by their structural formulas. An arrow is used to indicate the reaction, with the formulas for the starting materials on the left and those of the products on the right. In the structural formulas, the atoms of the elements are represented by their chemical symbols (C for carbon, H for hydrogen, O for oxygen), and the numbers of the atoms in particular groups are designated by numeral subscripts. The chemical bonds of greatest interest are represented by short lines between the symbols of the atoms connected by the bonds. Important to this reaction is an equilibrium involving the cleavage of the water molecules into positively and negatively charged particles (ions), as follows: In this equation, the numeral in front of the symbol for the water molecule indicates the number of molecules involved in the reaction. The composite arrow indicates that the reaction can proceed in either direction, starting material being converted to products and vice versa. In practice, both reactions occur together and a balance, or equilibrium, of starting materials and products is set up. The significance of this equilibrium for the hydrolysis of ethyl acetate is that any of the three entities (water molecules, hydronium, or hydroxide ions) may be involved in the reaction, and the mechanism is not known until it is established which of these is the actual participant. This often can be established if it is possible to determine the relative amounts of the three in the reaction medium and if it can be shown that the rate of the reaction depends upon the amount (or concentration) of one of them. Under certain conditions the hydrolysis of ethyl acetate is found to involve water molecules (as shown in the equation above); in other cases, hydroxide ion is involved. Mechanisms and kinetics Chemical kinetics Chemical reactions are explained in terms of the atomic structure of matter and of the energy changes that can take place in atoms, bonding them into molecules or breaking up molecules to free or to rebond atoms into different molecules. Chemical reactions may be slow or fast, complicated or relatively simple, and the modes by which they proceed are the subject of chemical kinetics. Preliminary considerations In a broader view, molecules have a tendency to react with each other to produce other kinds of molecules on condition that the product molecules have less free energy than the reacting molecules had; i.e., the products lie at a lower level of free energy than the reactants. An analogy is that of water always flowing downhill since, in the Earth's gravitational field, water at a higher level gives up energy as it falls to a lower level. Thus, all reactions for which the free energy, symbolized by DG, is negative have a tendency to occur. Accordingly, measurements of energy change suffice to indicate whether there will be a tendency for particular products to form when specific reactants are brought together under variable conditions. How fast a reaction will occur, however, depends on what sort of channels are available for the system to traverse in passing from the higher to the lower free energy state, and a study of the nature and effectiveness of these channels for reaction is best begun with a generalized symbolization of a reaction. In an actual reaction in which molecules of a substance A react with molecules of a substance B to form the molecules of substances C and D, usually the reverse reaction is also possible; that is, molecules of C and D react under suitable conditions to form molecules of A and B. If A, B, C, and D represent single molecules of those substances, a chemical equation may be written for the reaction: A + B C + D, the half arrows indicating that the reaction is reversible. Of course, if molecules A and B are to react they must first collide, and certain other conditions must also obtain. The number of collisions will be proportional to the concentrations of A and B molecules; if the amount of A is doubled, twice as many encounters will take place. The same is true if the amount of B is doubled. Thus, a change in the concentrations of the substances will change the number of collisions, and the more numerous these collisions are per second, the faster the reaction will proceed (assuming that other determining factors, such as temperature, remain constant). Thus, the rate of the reaction, or its speed, can be calculated by simply multiplying the concentrations. If concentration is signified by bracketing the symbols for the molecules, [A], [B], [C], and [D], and if vf symbolizes the forward rate, or the speed of reaction to the right in the equation, then this rate equals the product of the concentrations and of a proportionality factor, kf, also called the specific reaction-rate constant: vf = kf. Similarly, the backward rate is given by the equation vb = kb. When the two rates are the same-i.e., when vf = vb-the reaction, or, more correctly, the dynamic system of the two reactions, is said to be in equilibrium and kf = kb. Rearranging the equation, a ratio is obtained equal to a constant K, called the equilibrium constant; i.e., the product of the concentrations of the products divided by the product of the concentrations of the reactants is equal to the forward reaction-rate constant divided by the backward reaction-rate constant: This is known as the reaction-equilibrium equation. The equilibrium constant K, in view of the energy changes necessary for a reaction to proceed, further relates to the free energy, DG, which in turn relates to the forward and backward free energy of activation. The energy of activation is a measure of the energy that the reacting molecules must have in order to react, and it is specific for each reaction. Molecules may collide forever without reacting, and all the familiar solids, liquids, and gases consist of molecules in intimate contact but not reacting. Activation energy for the burning of paper, for example, can be provided to the paper molecules by the flame of a match, itself stable until the heat from friction on the head provides enough energy to activate the chemical reaction of the ingredients. The activation energies, therefore, of the forward (DGf) and backward (DGb) reactions are related to the free energy by the equation DG = DGf - DGb. Natural scientists have been impressed by the growth of crystals and of living things for centuries, but the first recorded, quantitative measurement of the rate of a chemical reaction was made in 1850, when it was noted that the plane of a beam of polarized light, passed through a solution of sucrose, was rotated by the solution and that the amount of rotation changed with time. This indicated that a chemical reaction was changing the sucrose into a new species of molecule (which rotated the polarized light to a different extent). The change in the concentration of sucrose can be measured and related to the passage of time to yield a specific rate constant for the reaction. It was found that the rate of any reaction depends not simply on the concentration of the reactants but on these concentrations raised to various powers. This principle is known as the law of mass action. It was pointed out by the Norwegian chemists C.M. Guldberg and Peter Waage around 1864-67 and precisely stated by a Dutch physical chemist, Jacobus Henricus van't Hoff, in 1877. In 1889 Svante Arrhenius, a Swedish chemist, published his conclusions that molecules must get into an active state before they can become reactive (i.e., their energies must be raised to activation levels before they will react). Van't Hoff had already obtained the result for the equilibrium constant, K, by calculations involving the energy of reaction, the absolute temperature, and the gas constant (derived from energy relationships observed in the behaviour of gases). Accordingly, Arrhenius was able to formulate an equation (named after him) for the specific rate constant of a chemical reaction that included the energy of activation.